Difference Between Vapor and Gas

Difference Between Vapor and Gas - Properties, FAQs

Edited By Team Careers360 | Updated on Jul 02, 2025 04:57 PM IST

To minimize the confusion, it's critical to understand the major differences between vapour and gas before using either term. Solid, Liquid, and Gas are the three states in which matter can be found to exist. Solids are predefined in terms of their shape and volume, unlike liquids (for example Ice). Water and other liquids have a specific volume, but depending on the container, they can take on different shapes.

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Vapour
Measuring Vapour
Examples of Gases and Vapour
Gases
Properties of Gases

The gaseous state is the third and final state, and it has neither a definite shape nor volume. Water vapour and oxygen are examples of gases. Matter can be transformed from one state to another by adjusting the temperature and pressure. When you lower the temperature or increase the pressure on a substance, the molecules get closer together and pack in more tightly.

This procedure permits the phase shifts from gas to liquids and from liquids to solid. The molecules move apart as the temperature and pressure are raised and lowered, and this can lead to phase transitions. from solids to liquids and also from liquids to gases. When normal conditions change, a material can skip a phase and proceed directly to the next.

Sublimation is the process by which a solid can be transformed directly into a gas without passing through an intermediate liquid stage. Dry ice, ammonium chloride, camphor, and iodine are a few examples of these substances. This article clears the questions like what is vapour, is water vapour a gas, what is the difference between gas and vapour etc.

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Vapour

Do you know what is meaning of vaporous? Let's look at what is vapour. Boiling a liquid produces vapour, which is created by a process called evaporation. Vapour definition is given as a state between the liquid and gas phases. When a liquid or solid is at equilibrium with it, it can coexist. Examples include water vapour and mercury vapour. Boiling or evaporation creates water in a gaseous state, which is referred to as water vapour.

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An aerosol is a mixture of gas and solid or liquid particles suspended in the gas. Water vapour has a greater greenhouse effect because it emits and absorbs infrared radiation over a wider range of wavelengths. In contrast to other greenhouse gases, water vapour makes up a disproportionately big portion of the earth's atmosphere. The amount of water vapour in the atmosphere changes constantly, making it impossible to get an accurate reading.

Measuring Vapour

Because it is in the gaseous state the partial pressure of the gases is used to calculate the amount of vapour present. In a gravitational field, vapours like normal atmospheric gases follow the Barometer Formula.

Examples of Gases and Vapour

As we know the most common example of vapour is water vapour (gas phase water) a room temperature and one atmosphere of pressure. And the second is gas; we know the most common example of gas is air (the air we breathe is gas). It can also be considered a mixture of many gases, such as oxygen, nitrogen, carbon dioxide, etc.

Gases

Gas is one of the four states of matter, and one of its distinguishing characteristics is that it takes up all of the available space, independent of its shape or volume. This property is due to the molecules having relatively low intermolecular attraction. Solid, liquid, and plasma are the other states of matter.

Solids have a fixed shape and volume, while liquids have a fixed volume but lack the fixed shape attribute; they take on the shape of the container they are poured into. Due to the weak interaction between the molecules, the gaseous molecules move constantly and independently. Due to the continual motion of the gaseous molecules, the gas can fill any size container. Compounds are categorised as gases even if they are still in a gaseous state. Carbon dioxide, for example, is a gas because it continues to be a gas at room temperature.

It is possible to create compressible fluid by compressing a gas's molecules, which are constantly moving past one another. Fixed gases are those gases that, regardless of temperature, cannot solidify or liquid is. Gases can't be seen with the human eye since their molecules are so far apart. At room temperature, gases such as carbon dioxide exist in gaseous form.

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Properties of Gases

They're easily compacted, which makes them useful. Intermolecular attraction keeps gas molecules apart in a steady state. Because of this, gases can be compressed with little effort, and this property is employed to compress gases in science, medicine, and engineering. The increased pressure and decreased temperature are used in compression.

Spray bottles, which compress the gases inside to make it easier to spray perfume when needed, are an excellent illustration of how air compression is used. No matter how big or small they are, they may all fit into the container. When you fill a balloon with air, the molecules are distributed evenly across the container. Gravity, on the other hand, limits the amount of liquid that may be put into a container.

Liquids in containers are shaped by gravity, whereas gases are less thick and can spread evenly. They take up more room than either solids or liquids. Oxygen and nitrogen are the most prevalent gases in the atmosphere (78 percent Nitrogen and 21 percent Oxygen). It contains a wide range of different gases, all in quite small concentrations. Difference between gases and vapors/ difference between steam and vapour:

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What is the Difference Between Gas And Vapour?

For the most part, a vapour phase is made up of two separate substances, one at ambient temperature and the other at a specific thermodynamic range. As a result, these are the primary differences between Vapor and Gas. This is the main difference between gas and vapour.
Vapours are usually solid or liquid, but they can become gaseous when exposed to certain conditions. It's not a physical property. Under normal circumstances, gases are in a gaseous state (at room temperature and one atmospheric pressure). A gaseous state of matter is one in which nothing exists. This is another difference between vapour and gas.
Whenever a vapour goes from liquid to vapour or from solid to vapour, it undergoes a phase change. The phase of a gas does not change. Any phase shift transforms a gas into a liquid. This is also a difference between gas and vapour.
The natural state of vapour is not gaseous, but might be solid or liquid. But the natural form of gas is in the gaseous state. This is another important difference between gas and vapour.
Under normal circumstances, the vapour is a substance that is both gaseous and liquid in nature. In normal circumstances, gas has a single thermodynamic state. This is another difference between gas and vapour.
The molecules and atoms that make up vapour move at random. Atoms and molecules in gases travel in random directions. This is also a difference between gas and vapour.

gas and vapour

In this way, gas and vapour difference is made.

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Frequently Asked Questions (FAQs)

1. What is the difference between a gas and vapour?

At normal temperature, vapour is a mixture of two or more distinct phases: the liquid and gaseous phases. At room temperature, gas normally has a single thermodynamic condition.When vapour is examined under a microscope, it reveals a clump of particles with no discernible shape.When gas is viewed under a microscope, it lacks a distinct shape.The molecules and atoms in a vapour are a random collection of particles travelling about in space.The molecules and atoms that make up gas are also made up of unorganised chaos.Contrary to gases, vapour is not a state of substance. The state of matter is gaseous.We are constantly surrounded by water vapour, even when the temperature is below the boiling point.Most gases arise when a substance reaches a temperature or pressure that is greater than or equal to its critical value.

2. What is vapour in chemistry?

It is the gaseous state of water.

3. What are the characteristics of vapour?

Under the critical temperature, a substance can exist in two states simultaneously: solid or liquid. This is called a vapour phase. When the vapour and either the solid or liquid phase are closely associated, both phases are in balance. The term "gas" often refers to a fluid phase that is easily compressed. Cloud formation is linked to vapour condensation. Vapor molecules have three types of motion: rotation, translation, and vibration.

4. What is vapour pressure?

When discussing vapour pressure, scientists refer to it as the pressure that solids or liquids exert when heated to a specific temperature. The vapour pressure is described by Raoult's law, which states that the partial pressure of any component is equal to the product of the pure component's vapour pressure and the mole fraction of the mixture.

5. What are the properties of gases?

It's the condition in which all of the particles are a long way apart. It's one of the three fundamental states of matter. When it comes to gases, the intermolecular distances between molecules are enormous and can be compressed with ease. It is impossible to measure the attractive forces that exist between the gas particles. The substances that exist as gaseous mixtures have neither a specific structure nor volume. They take up the entire volume if they're contained. They next apply a set amount of pressure to the container's walls.

6. How does the presence of a non-condensable gas affect the behavior of a vapor?

The presence of a non-condensable gas can significantly affect vapor behavior. It can lower the partial pressure of the vapor, affecting its condensation temperature and rate. This principle is important in processes like steam distillation, where a non-condensable gas can help volatilize substances at lower temperatures than their normal boiling points.

7. How do vapors and gases differ in their contribution to atmospheric phenomena?

Vapors, particularly water vapor, play a crucial role in atmospheric phenomena like cloud formation, fog, and precipitation. Gases like nitrogen and oxygen contribute to overall atmospheric composition and pressure. The interplay between vapors (e.g., water vapor) and gases in the atmosphere is key to understanding weather and climate patterns.

8. How does the concept of fugacity apply to vapors and gases?

Fugacity is a measure of the tendency of a substance to escape from a phase, often used to describe real gases and vapors that deviate from ideal behavior. For ideal gases, fugacity equals pressure. For vapors and real gases, fugacity accounts for non-ideal behavior, providing a more accurate description of their thermodynamic properties, especially at high pressures or near critical points.

9. What is the significance of Avogadro's law for understanding gases, and how does it apply to vapors?

Avogadro's law states that equal volumes of gases at the same temperature and pressure contain the same number of molecules. This law applies more accurately to gases than to vapors. Vapors, being closer to their condensation point, may deviate from this law due to stronger intermolecular interactions and non-ideal behavior.

10. How do vapors and gases differ in their heat capacity behaviors?

The heat capacity of gases is generally lower and more constant over a wide temperature range compared to vapors. Vapors, especially near their condensation point, can have more complex heat capacity behaviors due to changes in intermolecular interactions as they approach the liquid state. This difference affects how they absorb and release heat in various processes.

11. How does the ideal gas law apply differently to vapors and gases?

The ideal gas law (PV = nRT) applies more accurately to gases, especially at high temperatures and low pressures. Vapors deviate more from ideal behavior due to stronger intermolecular forces and being closer to their condensation point. However, as vapors are heated and move further from their condensation point, they begin to behave more like ideal gases.

12. How does the concept of saturation apply differently to vapors and gases?

Saturation applies primarily to vapors. A saturated vapor is in equilibrium with its liquid or solid phase, and its pressure is the maximum vapor pressure at that temperature. Gases, being above their critical temperature, cannot be saturated in the same way, as they cannot condense into a liquid regardless of pressure.

13. What role do intermolecular forces play in distinguishing vapors from gases?

Intermolecular forces are stronger in vapors compared to gases. In vapors, these forces are significant enough to allow relatively easy condensation back to the liquid state. In gases, intermolecular forces are much weaker, allowing the molecules to move more freely and making it harder to condense them into liquids without extreme conditions.

14. What is the significance of Dalton's law of partial pressures for mixtures of vapors and gases?

Dalton's law states that the total pressure of a mixture of gases is equal to the sum of the partial pressures of each gas. This law applies to both vapors and gases in a mixture, allowing us to calculate the contribution of each component to the total pressure, regardless of whether it's a vapor or a gas.

15. How does the density of a vapor compare to that of a gas?

Vapors generally have a higher density than gases at the same temperature and pressure. This is because vapor molecules are closer together due to stronger intermolecular attractions, while gas molecules are more spread out. However, as temperature increases, the density difference becomes less pronounced.

16. What is vapor pressure, and how does it relate to the concept of vapor?

Vapor pressure is the pressure exerted by a vapor in equilibrium with its liquid or solid phase at a given temperature. It's a measure of a liquid's tendency to vaporize. As temperature increases, vapor pressure increases because more molecules have enough energy to overcome intermolecular forces and enter the vapor phase.

17. What is the significance of Raoult's law in understanding vapor behavior?

Raoult's law describes the vapor pressure of an ideal solution, stating that the partial vapor pressure of each component is proportional to its mole fraction in the solution. This law is crucial for understanding the behavior of mixed liquids and their vapors, such as in distillation processes or the formation of azeotropes.

18. What is the significance of the triple point in understanding vapors?

The triple point is the unique combination of temperature and pressure at which a substance can exist simultaneously as a solid, liquid, and vapor in equilibrium. It's significant for understanding vapors because it represents the lowest pressure at which a liquid can exist, below which only the solid and vapor phases are possible.

19. What is the relationship between vapor pressure and atmospheric pressure?

When a liquid's vapor pressure equals atmospheric pressure, the liquid boils. At normal atmospheric pressure (1 atm or 101.325 kPa), this occurs at the normal boiling point. If atmospheric pressure is lower (e.g., at high altitudes), liquids will boil at lower temperatures because their vapor pressure equals the atmospheric pressure at a lower temperature.

20. What is the difference between saturated and unsaturated vapors?

A saturated vapor is in equilibrium with its liquid phase at a given temperature, exerting its maximum vapor pressure. An unsaturated vapor has a pressure lower than its saturation vapor pressure at that temperature. Unsaturated vapors can accept more molecules into the vapor phase, while saturated vapors are at maximum capacity.

21. How does temperature affect the behavior of vapors and gases?

Temperature affects both vapors and gases by increasing molecular kinetic energy. For vapors, increasing temperature can cause them to behave more like ideal gases. For gases, temperature increase leads to higher pressure if volume is constant, or expansion if pressure is constant, following Gay-Lussac's and Charles's laws, respectively.

22. What is the critical temperature, and how does it relate to vapors and gases?

The critical temperature is the highest temperature at which a substance can exist as a liquid. Above this temperature, the substance can only exist as a gas, regardless of pressure. Vapors can be liquefied below their critical temperature, while gases above their critical temperature cannot be liquefied by pressure alone.

23. How do vapors and gases differ in terms of compressibility?

Gases are generally more compressible than vapors. This is because gas molecules are farther apart and have weaker intermolecular forces, allowing them to be easily pushed closer together. Vapors, being closer to their liquid state, have stronger intermolecular forces and are less compressible.

24. What is the main difference between a vapor and a gas?

The main difference is that a vapor is the gaseous state of a substance that is normally a liquid or solid at room temperature, while a gas is a substance that exists in the gaseous state under normal conditions. Vapors can be easily condensed back to their liquid or solid form, whereas gases require extreme cooling or pressure to liquefy.

25. Why can vapors be liquefied more easily than gases?

Vapors can be liquefied more easily because they are closer to their condensation point. The molecules in vapors have weaker intermolecular forces and higher potential energy compared to their liquid state, making it easier to overcome these forces and return to the liquid form with less energy input.

26. What is the relationship between boiling point and the formation of vapors?

The boiling point is the temperature at which a liquid's vapor pressure equals the atmospheric pressure, causing the liquid to vaporize rapidly. At temperatures below the boiling point, vapors still form (through evaporation), but at a slower rate. The lower the boiling point of a substance, the more readily it forms vapors at room temperature.

27. What is the difference between evaporation and vaporization?

Evaporation is the process by which molecules escape from the surface of a liquid to form a vapor at any temperature, occurring only at the liquid-air interface. Vaporization, on the other hand, is the general term for the change from liquid to gas and includes both evaporation and boiling (rapid vaporization throughout the liquid).

28. What is supersaturation, and how does it relate to vapors?

Supersaturation is a state where a solution contains more dissolved material than could be dissolved by the solvent under normal circumstances. In the context of vapors, a supersaturated vapor contains more gaseous molecules than the equilibrium amount at a given temperature and pressure, making it unstable and prone to rapid condensation.

29. How does pressure affect the behavior of vapors differently from gases?

Pressure has a more pronounced effect on vapors than on gases. Increasing pressure on a vapor can cause it to condense into a liquid more easily, especially near its critical point. Gases, being further from their condensation point, require much higher pressures to liquefy and behave more predictably under pressure changes, following Boyle's law more closely.

30. How does the presence of a vapor affect the properties of a gas mixture?

The presence of a vapor in a gas mixture can significantly alter its properties. Vapors can condense more easily than gases, potentially leading to liquid formation under certain conditions. They also contribute to the total pressure of the mixture according to Dalton's law, and can affect properties like density, viscosity, and heat capacity of the overall mixture.

31. How does the kinetic molecular theory explain the behavior of vapors and gases?

The kinetic molecular theory explains that both vapors and gases consist of particles in constant, random motion. However, in vapors, the particles have stronger intermolecular attractions and less kinetic energy compared to gases. This results in vapors being more easily condensed and deviating more from ideal gas behavior.

32. How do vapors and gases differ in their response to cooling?

When cooled, vapors more readily condense into liquids or solids as they approach their condensation point. Gases, being further from their condensation point, require more extreme cooling to liquefy. The cooling process reduces the kinetic energy of the particles, making it easier for intermolecular forces to dominate in vapors.

33. How does the concept of mean free path differ for vapors and gases?

The mean free path, the average distance a particle travels between collisions, is generally shorter for vapors than for gases. This is because vapor molecules are closer together and have stronger intermolecular attractions, leading to more frequent collisions. In gases, molecules are farther apart, resulting in a longer mean free path.

34. How do vapors and gases differ in their behavior under isothermal compression?

Under isothermal compression, gases generally follow Boyle's law more closely, with pressure increasing inversely proportional to volume. Vapors, especially near their condensation point, deviate from this behavior. As pressure increases, vapors may begin to condense, leading to a more complex pressure-volume relationship compared to gases.

35. What is the relationship between vapor pressure and intermolecular forces?

Vapor pressure is inversely related to the strength of intermolecular forces. Substances with stronger intermolecular forces (like hydrogen bonding in water) have lower vapor pressures because more energy is required for molecules to escape the liquid phase. Conversely, substances with weaker intermolecular forces have higher vapor pressures.

36. What is the relationship between vapor density and molecular weight?

Vapor density is directly related to molecular weight. Heavier molecules form denser vapors compared to lighter molecules under the same conditions. This relationship is important in understanding vapor behavior in mixtures and in processes like distillation, where differences in vapor density can be used to separate components.

37. What is the difference between vapor-liquid equilibrium and gas-liquid equilibrium?

Vapor-liquid equilibrium refers to the state where a vapor is in equilibrium with its own liquid phase, with equal rates of vaporization and condensation. Gas-liquid equilibrium, on the other hand, typically refers to the equilibrium between a dissolved gas and its gaseous form above a liquid, governed by Henry's law. The key difference lies in the nature of the gaseous phase and its relationship to the liquid.

38. How does the concept of critical pressure relate to vapors and gases?

Critical pressure is the pressure required to liquefy a gas at its critical temperature. Above the critical pressure and temperature, the distinction between liquid and gas phases disappears. For vapors, the critical pressure represents the maximum pressure at which the substance can exist as a vapor. Gases above their critical pressure and temperature are often referred to as supercritical fluids.

39. What is the significance of the Clausius-Clapeyron equation in understanding vapor behavior?

The Clausius-Clapeyron equation describes how the vapor pressure of a liquid changes with temperature. It's crucial for understanding phase transitions, especially vaporization. The equation relates the slope of the vapor pressure curve to the enthalpy of vaporization and temperature, allowing predictions of vapor pressure at different temperatures and calculation of enthalpies of vaporization.

40. How do vapors and gases differ in their behavior during adiabatic processes?

During adiabatic processes (no heat exchange with surroundings), gases typically follow the adiabatic gas law more closely, with temperature and pressure changes related by the heat capacity ratio. Vapors, especially near condensation, may deviate from this behavior due to phase changes. The adiabatic expansion of a vapor can lead to condensation, while gases remain in the gaseous state.

41. What is the role of vapor pressure in determining the volatility of a substance?

Vapor pressure is a key determinant of a substance's volatility. Substances with higher vapor pressures at a given temperature are more volatile, meaning they evaporate more readily. This concept is crucial in understanding evaporation rates, boiling points, and the behavior of mixtures, including the formation of azeotropes and the principles behind distillation.

42. How does the presence of dissolved solids affect the vapor pressure of a liquid?

Dissolved solids generally lower the vapor pressure of a liquid, a phenomenon known as vapor pressure lowering. This occurs because the presence of solute particles reduces the number of solvent molecules at the surface available for evaporation. This principle is the basis for colligative properties like boiling point elevation and is crucial in understanding solutions and their vapor behaviors.

43. What is the difference between an ideal vapor and a real vapor?

An ideal vapor follows the ideal gas law perfectly, assuming no intermolecular forces and point-like molecules. Real vapors deviate from this behavior, especially near their condensation point, due to intermolecular forces and the finite size of molecules. The deviation becomes more pronounced at higher pressures and lower temperatures, where real vapors may condense while an ideal vapor would not.

44. How does the concept of partial pressure apply differently to vapors in a mixture compared to gases?

While partial pressure applies to both vapors and gases in mixtures (following Dalton's law), its implications can differ. For gas mixtures, partial pressures often behave additively. In vapor mixtures, especially those close to condensation, interactions between different vapor species can lead to non-ideal behavior, affecting the relationship between partial pressures and overall mixture properties.

45. What is the significance of the vapor pressure curve in a phase diagram?

The vapor pressure curve in a phase diagram represents the boundary between the liquid and vapor phases of a substance. It shows how vapor pressure changes with temperature and helps predict phase transitions. The curve ends at the critical point, beyond which the distinction between liquid and vapor phases disappears. Understanding this curve is crucial for predicting a substance's phase at various temperatures and pressures.

46. How do vapors and gases differ in their behavior during Joule-Thomson expansion?

During Joule-Thomson expansion (an isenthalpic process), gases typically cool upon expansion, with the extent depending on their initial temperature and pressure. Vapors, especially near their condensation point, may exhibit more complex behavior. They can experience more significant cooling or even condensation during expansion, depending on their initial state relative to their critical point and inversion temperature.

47. What is the relationship between vapor pressure and boiling point elevation in solutions?

Boiling point elevation in solutions is directly related to vapor pressure lowering. When a non-volatile solute is added to a solvent, it lowers the solvent's vapor pressure. As a result, a higher temperature is required for the solution's vapor pressure to equal atmospheric pressure, leading to a higher boiling point. This relationship is quantified by Raoult's law and is a key concept in understanding colligative properties.

48. How does the behavior of vapors near their critical point differ from that of gases?

Near the critical point, vapors exhibit highly non-ideal behavior. They become highly compressible, and their properties (like density) can change dramatically with small changes in temperature or pressure. Gases, unless near their own critical points, maintain more consistent and predictable behavior. The distinction between liquid and vapor phases becomes blurred for a substance near its critical point, unlike for gases well above their critical temperatures